Quantitative Assessment of the Connection between Steric Hindrance

May 27, 2014 - ... Assessment of the Connection between Steric Hindrance and Electronic Coupling in 2,5-Bis(alkoxy)benzene-Based Mixed-Valence Dimers...
0 downloads 0 Views 1MB Size
Download PDF
Article pubs.acs.org/JPCC

Quantitative Assessment of the Connection between Steric Hindrance and Electronic Coupling in 2,5-Bis(alkoxy)benzene-Based Mixed-Valence Dimers Angela M. Bischof,† Shaopeng Zhang,‡ Tara Y. Meyer,‡ and Benjamin J. Lear*,† †

Department of Chemistry, The Pennsylvania State University, University Park, Pennsylvania 16802, United States Department of Chemistry, University of Pittsburgh, 219 Parkman Avenue, Pittsburgh, Pennsylvania 15260, United States

‡

S Supporting Information *

ABSTRACT: The effect of the bridging ligand on electronic delocalization was examined in a series p-bis(alkoxy)benzene dimers relevant to conducting polymers used for organic devices. Using spectroscopic methods, the degree of delocalization for an ethylene-bridged p-bis(alkoxy)benzene dimer was determined and compared to the electronic coupling for directly coupled and phenylene-bridged p-bis(alkoxy)benzene dimers reported previously. Despite a significant increase in distance (53%) between the redox-active sites, the ethylene-bridged compound exhibited a higher electronic coupling than either of the others previously reported. The increased coupling can be attributed to the lower rotational barrier to planarization for the ethylene-bridged dimer. This result highlights the need to minimize both sterics and distance between redox active sites in molecular systems designed for promoting electron mobility and provides quantitative evidence that an optimal balance between these parameters can be achieved.

■

INTRODUCTION Charge mobility is a key concern for the design of conjugated organic polymers, which continue to be relevant for emerging technologies such as organic light-emitting diodes (LEDs), field effect transistors, and photovoltaic devices. Maximizing charge mobility in these films involves a balance between maximizing electronic coupling between the redox units of the polymer (which requires placing redox units in close proximity) and minimizing rotational steric hindrance between the repeating units of the polymer (which requires spatial separation of redox units). Proper balancing of these parameters is critical for controlling both intrachain and interchain electron transfer which together determine the overall charge mobility of the film. With regard to the overall electron mobility within polymers, we can consider the anticipated impact of the above parameters: electronic coupling and steric hindrance. The electronic coupling between redox units affects the electron mobility by controlling the conjugation length of the polymer, with electronic coupling increasing as a function of orbital overlap and exponentially decreasing with distance between repeat units.1,2 As the effective conjugation length increases, both the mobility along the polymer chain and the rate of interchain hopping are expected to increase. The latter of these effects stems from an increase in the overlap between electron wave functions on separate chains (a direct result of increased conjugation length) as well as a reduction in the reorganization energy for electron transfer that itself results from increased distribution of the charge among the bonds of the polymer. Turning to steric considerations, rotational sterics affect charge © 2014 American Chemical Society

mobility by controlling the planarity of the polymer backbone which, in turn, affects both electronic coupling along an individual chain and molecular packing between chains. Molecular packing is particularly important for controlling the rate of interchain hopping, which often imposes the ultimate limit on charge mobility, and can be affected by the presence and identity of side chains3 and the manner in which the monomers are joined together.4−6 However, the simplest way to reduce steric hindrance is to increase the distance between redox units, and this places the desire for increased planarity in direct conflict with the desire to maximize electronic coupling. Thus, the design of future conjugated polymers will benefit from a quantitative understanding of the tension between the competing needs to maximize both electronic coupling and planarity. Our approach to providing this quantitative understanding is to use dimers of redox active units common to conjugated polymers as model systems. The validity of this approach is supported by recent computational studies that indicate many of the relevant properties of polymeric systems are already emergent when considering such dimeric models.7,8 Due to their inherent redox activity, these dimeric models can be considered mixed-valence compounds, a class of compounds long used to investigate fundamental electron transfer problems, and this perspective allows for many of the tools of mixed-valence chemistry to be applied to these systems. One Received: May 18, 2014 Revised: May 27, 2014 Published: May 27, 2014 12693

dx.doi.org/10.1021/jp504887s | J. Phys. Chem. C 2014, 118, 12693−12699

The Journal of Physical Chemistry C

Article

Figure 1. Two-state model showing potential energy surfaces for Robin and Day Classes I (left panel), II (center panel), and III (right panel). The green curves shown in the center and right panels are the same diabatic curves found in the left panel.

the length of the bridge (Chart 1).11 An additional point of interest is that, to the best of our knowledge, 2 is the first mixed-valence compound in which the bridge is clearly a single ethylene unit. Though there are prior examples of mixedvalence compounds that contain an ethylene moiety in the bridge,12,13 ambiguity concerning the dissection of the molecule into its redox active centers and bridge renders clear discourse on the nature of the bridge difficult. For this reason, we believe that 2 is the most definitive example of a mixed-valence molecule in which two redox sites are bridged by only an ethylene linker. Because electronic coupling decreases exponentially with distance,1,2 it seems reasonable to hypothesize that 2•+ will be less strongly coupled than 1•+ but more strongly coupled than 3•+. However, we find that 2•+ exhibits greater electronic coupling than 1•+ or 3•+, which we attribute to reduction of steric hindrance in 2 versus 1. Below, we discuss this result, quantify the cost of steric hindrance in terms of electronic coupling, and emphasize the importance of considering not only the properties of the redox active repeat unit but also the manner of connecting these units.

such tool is the two-state Marcus−Hush model, which dominates the theoretical interpretation of the behavior of mixed-valence compounds and allows for quantitative discussion of the electronic coupling present in these dimers.1 In this model, the two electronic distributions are represented by two diabatic potential energy surfaces (Figure 1). Electronic coupling between these states generates lower and upper adiabatic potential energy surfaces. The extent of this coupling determines the shape of the lower potential energy surface, which is used as the basis for classification of mixed-valence compounds within the Robin−Day scheme.9 Within this scheme molecules are assigned to one of three classes, according to the extent of their electronic coupling and the localization of the unpaired electron: Class I (uncoupled, localized electron); Class II (coupled, localized electron); and Class III (coupled, delocalized electron). The magnitude of electronic coupling for these molecules can be determined through the electrochemistry and spectroscopy of the molecules. This provides a ready means by which to quantify the electronic coupling, and herein, we quantify the connection between the extent of electronic coupling and the steric barrier to planarization of the redox active units. Our focus is on the electronic properties of (E)-1,2-bis(2,5bis(hexyloxy)-4-methylphenyl)ethene (2). This molecule can be considered the dimeric analogue of 2,5-bis(alkoxy)substituted poly(p-phenylenevinylene) (PPV) polymers.10 In addition, similar molecules, 1 and 3 (see Chart 1), have also

■

EXPERIMENTAL AND THEORETICAL METHODS Synthesis. Synthesis of 2 followed the route presented in Scheme 1. Compound D-CHO was prepared as previously reported.10 KOtBu (1.0 M in THF) was purchased from Aldrich and dispensed using air-sensitive techniques. NaBH4 was stored in a desiccator over anhydrous CaSO4. LiCl was dried at 120 °C for at least 24 h. Reagent grade THF was used for reactions; notably the HWE reactions used reagent grade THF. All other reagents and solvents were used as received. All reactions were carried out under N2 atmosphere. Column chromatography was carried out on standard grade silica gel (60 Å pore size, 40−63 μm particle size), which was purchased and used as received. Hexanes, dichloromethane, and ethyl acetate used for column chromatography were purchased and used as received. HPLC-grade solvents were purchased from Sigma-Aldrich and were further purified using an Innovative Technology Pure Solv system for all spectroscopic measurements. Dimethyl 2,5-Bis(hexyloxy)-4-methylbenzylphosphonate (P-D). P-D was synthesized by a slight modification of Jorgensen’s method14 in three steps: (1) D-CHO (3.00 g, 9.36 mmol) was dissolved in THF (50 mL) and ethanol (25 mL) and treated with sodium borohydride (2 g, excess) for 1 h. The solvents were removed under vacuum, and the white solid was redissolved in ether (50 mL) and washed with water (50 mL). The organic phase was dried over magnesium sulfate, and the solvent was removed under vacuum. (2) The white solid was dissolved in THF (50 mL). Carbon tetrabromide (3.41 g,

Chart 1. Structures of Mixed-Valence Compounds

a

The properties of 1 and 3 have been reported by Kochi and coworkers (ref 11).

been investigated using mixed-valence approaches,11 which allows us to discuss the distance dependence of electronic coupling and rotational sterics as a function of distance between redox sites for these compounds. Compound 1 is the dimeric analogue of 2,5-bis(alkoxy)-substituted poly(p-phenylene) (PPP) polymers, and 3 is the dimeric analogue of a PPP coblock polymer. The key difference among these molecules is 12694

dx.doi.org/10.1021/jp504887s | J. Phys. Chem. C 2014, 118, 12693−12699

The Journal of Physical Chemistry C

Article

Scheme 1. Synthesis Route of Compound 2 (DD)

methylene chloride in hexanes). Recrystallization from methanol and hexanes gave compound 2 as a yellow solid (158 mg, 0.26 mmol, 77%). 1H NMR (400 MHz, CDCl3) δ 7.38 (s, 2H), 7.07 (s, 2H), 6.71 (s, 2H), 3.96 (dd, J = 14.6, 6.5 Hz, 8H), 2.23 (s, 6H), 1.88−1.72 (m, 8H), 1.57−1.43 (m, 8H), 1.41−1.28 (m, 16H), 0.90 (q, J = 7.0 Hz, 12H). 13C NMR (126 MHz, CDCl3) δ 151.60 (s), 150.40 (s), 127.26 (s), 125.51 (s), 122.72 (s), 116.27 (s), 109.28 (s), 69.81 (s), 68.83 (s), 31.66 (d, J = 3.1 Hz), 29.56 (d, J = 2.2 Hz), 25.87 (t, J = 8.1 Hz), 22.66 (d, J = 1.9 Hz), 16.40 (s), 14.06 (d, J = 2.5 Hz). HRMS calculated for C40H65O4: 609.4882 g/mol. Found: 609.4883 g/ mol. The mixed-valence radical cation (2•+) was prepared by chemical oxidation of the neutral species using AgPF6 in dry dichloromethane under inert atmosphere. The dication species (22+) was prepared by chemical oxidation using NOPF6 in dry dichloromethane under inert atmosphere. Experimental Methods. 1H (400 MHz) and 13C (126 MHz) NMR spectra were recorded on Bruker spectrometers. Chemical shifts were referenced to residual 1H or 13C signals in deuterated solvents (7.27 and 77.0 ppm, respectively, for CDCl3 and 5.32 and 54.0 ppm, respectively, for CD2Cl2). Highresolution mass spectra (HRMS) were recorded on EIquadrupole or ESI-TOF instruments in the Mass Spectrometry Facility of the University of Pittsburgh. Cyclic voltammetry (CV) was performed using a Pine WaveNow Potentiostat. CVs were collected using a threeelectrode system consisting of a 3 mm Pt disk working electrode, a Pt wire counter electrode, and a nonaqueous Ag/ Ag+ quasi-reference electrode (QRE) at a scan rate of 100 mV/ s in 0.1 M NBu4PF6 in CH2Cl2 under a nitrogen atmosphere. Ferrocene was added as an internal standard to calibrate the Ag/Ag+ QRE. Infrared and NIR spectra were obtained using a PerkinElmer Spectrum 400 FT-IR/FT-NIR spectrometer. All infrared and

10.3 mmol, 1.1 equiv) and triphenylphosphine (2.70 g, 10.3 mmol, 1.1 equiv) were added to the solution, and the reaction mixture was allowed to stand overnight. The solvent was removed under vacuum, and the residue was dissolved in diethyl ether and filtered through a plug of silica. After the solvent was removed, the material was purified by column chromatography (5% ethyl acetate in hexanes) to give a yellow solid. (3) The yellow solid was suspended in trimethylphosphite (15 mL, excess) and was heated to reflux for 1.5 h, and the excess trimethylphosphite was removed under vacuum. The material was purified by column chromatography (25% ethyl acetate in hexanes) to compound P-D as a yellow oil (1.81 g, 4.37 mmol, 47%). 1H NMR (400 MHz, CD2Cl2) δ 6.80 (d, J = 2.5 Hz, 1H), 6.71 (s, 1H), 3.96−3.83 (m, 4H), 3.67 (s, 3H), 3.64 (s, 3H), 3.21 (s, 1H), 3.16 (s, 1H), 2.20 (d, J = 2.3 Hz, 3H), 1.85−1.69 (m, 4H), 1.57−1.42 (m, 4H), 1.41−1.29 (m, 8H), 0.92 (dd, J = 6.9, 5.8 Hz, 6H). 13C NMR (126 MHz, CDCl3) δ 151.00 (d, J = 3.5 Hz), 150.30 (d, J = 7.5 Hz), 126.56 (d, J = 4.0 Hz), 117.32 (d, J = 9.5 Hz), 115.06 (d, J = 2.9 Hz), 114.61 (d, J = 5.0 Hz), 69.17 (s), 68.83 (s), 52.65 (d, J = 6.6 Hz), 31.61 (d, J = 1.9 Hz), 29.48 (d, J = 12.1 Hz), 26.07 (s), 25.80 (d, J = 3.0 Hz), 24.96 (s), 22.62 (s), 16.29 (s), 14.01 (s). HRMS calculated for C22H40O5P: 415.2627 g/mol. Found: 415.2613 g/mol. (E)-1,2-Bis(2,5-bis(hexyloxy)-4-methylphenyl)ethene (2 = DD). D-CHO (0.108 g, 0.336 mmol), P-D (0.209 g, 0.504 mmol), and LiCl (0.0329 g, 0.773 mmol) were dissolved in THF (10 mL) and stirred under N2. KOtBu (1.0 M in THF, 0.77 mL, 0.77 mmol) was added dropwise. The reaction was allowed to stand overnight with stirring. Saturated NH4Cl water solution (10 mL) was added to the reaction mixture, and the reaction mixture was extracted with methylene chloride (3 × 10 mL). The combined organic phase was dried over magnesium sulfate, and the solvent was removed under vacuum. The material was first purified by column chromatography (30% 12695

dx.doi.org/10.1021/jp504887s | J. Phys. Chem. C 2014, 118, 12693−12699

The Journal of Physical Chemistry C

Article

equilibrium constant (Kcom) of 1.0 × 106 (eqs 1 and 2), indicating that the mixed-valence state is very stable.

NIR spectra were recorded using a Precision demountable liquid cell with KBr (IR) or CaF2 (NIR) windows and a 0.2 mm path length spacer. Samples for NIR spectroscopy were 2 × 10−4 M in dry dichloromethane. UV−visible spectra were recorded using an Agilent 8453 UV−visible spectrometer and a 1 cm quartz cuvette. All UV− vis samples were prepared at a concentration of 5 × 10−5 M in dry dichloromethane. UV−visible and NIR spectroelectrochemistry (SEC) were performed using a Pine WaveNow Potentiostat with a Pine Honeycomb electrode (Au working electrode and reference electrode) and a nonaqueous Ag/Ag+ reference electrode in a 1 mm path length quartz cuvette. UV− vis SEC was performed on an Agilent 8453 dispersive spectrometer, and NIR SEC was performed on a PerkinElmer Spectrum 400 FT-IR/FT-NIR spectrometer. UV−vis SEC samples were 1 × 10−4 M in 0.1 M Bu4NPF6 in dry dichloromethane, and NIR SEC samples were 2 × 10−3 M in 0.1 M Bu4NPF6 in dry dichloromethane. All SEC studies were performed by collecting spectra every minute during bulk electrolysis (BE). In both UV−vis and NIR SEC, BE experiments were performed by applying the potentials in the following order on a single solution: 0.8 V vs Ag/Ag+ to oxidize 20 to 2•+, 1.1 V vs Ag/Ag+ to oxidize 2•+ to 22+, 0.8 V vs Ag/ Ag+ to reduce 22+ to 2•+, and 0 V (UV−vis) or −0.1 V (NIR) to reduce 2•+ to 20. Raman spectra were recorded using a Renishaw inVia Microscope equipped with an integral microscope (Leica DM2500M). The excitation source was a 647 nm CrystaLaser CL-2000 diode pumped laser (70 mW, model DL647-070), and a 1200 L/mm grating was used, giving a resolution of 1.9 cm−1. Theoretical Methods. All structures were generated and viewed using GaussView 5.0.9.15 Structural optimization and energy calculations were performed using Gaussian0916 with density functional theory (DFT) B3LYP/6-31G. The barrier to rotation, ΔErot, was calculated for 1 and 2 by taking the singlepoint energy as a function of dihedral angle and finding the difference between the highest and lowest energy conformations. All calculations for 2 were done using methyl groups in place of the hexyl groups to reduce computational time.

K com

D−D2 + + D−D ←→ ⎯ 2D−D•+

(1)

ΔGcom = −F ΔE1/2 = −RT ln(Kcom)

(2)

This value of ΔE1/2 is larger than the value of ΔE1/2 reported by Kochi and co-workers for the stable mixed-valence cation 1•+ (0.29 V) and for the nearly delocalized cation, 3•+ (0.11 V), consistent with the idea that electronic coupling in 2•+ is greater than in 1•+ and 3•+. While one must be careful not to overinterpret values of ΔE1/2 because effects other than electronic coupling can control the magnitude of this parameter,17−19 in this case such analysis may be justified. In most cases, the dominant contribution to ΔE1/2 (other than electronic coupling) is electrostatic repulsion associated with adding a second charge spatially close to the first.17,18 As the distance between redox units in 2•+ is farther than in 1•+, however, it seems likely that the increase in ΔE1/2 for 2•+ (when compared to 1•+) is a result of increased couplingand that this increase in coupling should be greater than or equal to 0.16 V. While the electrochemistry gives qualitative indications of the degree of electronic coupling in 2•+, the electronic spectra can provide quantitative information. The UV−visible electronic absorption spectra of 2, 2•+, and 22+ (generated via chemical oxidation) are presented in Figure 2. The UV region is

■

Figure 2. UV−visible electronic absorption spectra of 2, 2•+, and 22+ in CH2Cl2. Both 2•+ and 22+ were generated by chemical oxidation.

RESULTS AND DISCUSSION The first indication that electronic coupling in 2 is greater than in 1 or 3 is found in the electrochemistry of the compounds. The cyclic voltammogram of 2 (see Supporting Information, Figure S2) shows two well-resolved reversible couples at 0.47 and 0.82 V vs ferrocene. The first of these events is associated with the oxidation of 2 to the mixed-valence state, 2•+, and the second to the generation of the dication. The fact that we only observe two couples indicates that there is only a single species present in the sample. This observation is important, as 2 could exist in either the E or Z configuration. Though we have prepared the pure E isomer, it is important to look for evidence of conversion to the Z isomer. Because the distance between the two redox centers would be different in the E and Z isomers, we would expect to observe a different degree of splitting between the redox events for these two isomers. Thus, the presence of only two redox couplings in the CV indicates that 2 is present in a single conformation, which we take to be the as-prepared E isomer. The separation of the standard oxidation potentials (ΔE1/2) is 0.35 V, which corresponds to a comproportionation

consistent with that of the monomeric prototype (1,4dimethoxy-2,5-dimethylbenzene) reported by Kochi and can be assigned to the p-alkoxybenzene subunit. In the visible region, the first oxidation of 2 results in the appearance of an intense absorption band at λmax = 502 nm with a low intensity shoulder at 460 nm and another weak band at λmax = 585 nm (see Table S1, Supporting Information, for extinction coefficients for all electronic absorptions). The behavior is also qualitatively similar to that observed by Kochi et al. when oxidizing the mononuclear analogue to the radical cation.11 This transition in the visible region was assigned to the interaction between the redox active site and the bridge, and we assume the same identity here.11 The blue-shift observed for 2•+ with respect to 1•+ then points to an increased energy gap between the redox active unit and the ethylene bridge. After the second oxidation, these peaks disappear, and new, weaker absorptions appear at λmax = 458, 486, 761, 929, and 994 nm. Previous studies do not thoroughly report the UV−visible absorptions for the dication species, and so, we have no prior 12696

dx.doi.org/10.1021/jp504887s | J. Phys. Chem. C 2014, 118, 12693−12699

The Journal of Physical Chemistry C

Article

= 1374 cm−1), indicating increased coupling in 2•+.12,21 A full description of how Δν was determined can be found in the Supporting Information. In total, however, the shape of the IVCT band supports the assignment of 2•+ as more strongly coupled than 1•+. Besides the shape, the intensity of IVCT bands is also sensitive to the degree of coupling, with the intensity increasing as a function of electronic coupling.21 We find that the IVCT band for 2•+ is more intense (ε = 1.5 × 104) than that observed for 1•+ (ε = 4.8 × 103).11 Again, this indicates that 2•+ is more strongly coupled than 1•+.22 At this point, we must mention that we are treating 2•+ as a fully delocalized (Class III) system (see below for further justification of this assignment), which makes the relative discussion of intensities less straightforward than for Class II systems.21 Nevertheless, as 1•+ and 2•+ are analogous systems, we feel that this comparison is justified. Finally, the position of the IVCT band allows for easy quantitative discussion of the electronic coupling in these systems using Marcus−Hush theory. Up to this point, all of our available evidence suggests that 2•+ is more strongly coupled than the Class III 1•+ cation. For this reason, we feel justified in assigning 2•+ as a fully delocalized (Class III) system. Furthermore, we will use the two-state model for mixed valency to compare our results for 1•+ to those reported by Kochi and co-workers for 2•+ and 3•+. Though the presence of possible vibronic features in the IVCT band calls into question the validity of the two-state model, this appraoch allows for direct comparison of 2, with 1 and 3. It is also worth noting that the electronic coupling for 3•+ was determined using the Hush equation for Class II mixed-valence species, which may underestimate the value of the electronic coupling for a Class II/III mixed-valence species, as shown by Kochi and coworkers.11 However, we feel the perceived usefulness of this treatment outweighs any subtle conceptual inaccuracies associated with using the two-state model. The treatment of 2•+ as a Class III system simplifies the extraction of the magnitude of the electronic coupling, and the electronic coupling (HAB) is easily obtained from the position (νmax) of of the lowest-energy IVCT band.

data with which to compare these results. However, based on UV−visible spectroelectrochemistry (Figure S2, Supporting Information), we are able to conclude that the complex electronic absorption spectra are due to electronic changes, not chemical changes, because the changes are reversible and the original spectra are recovered upon reduction. In the NIR region of the electronic spectrum, a broad asymmetric band appears upon oxidation from 2 to 2•+ by both chemical and electrochemical means. The results of the chemical oxidation are presented in Figure 3. The feature

Figure 3. NIR electronic absorption spectra of 2 (green) and 2•+ (black) with Gaussian deconvolution of 2•+ (red) and Gaussian fit (blue circles). Both 2•+, and 22+ were generated by chemical oxidation.

present in this spectrum is similar to that observed for similar molecules, where it was found to be unique to the full compound and was not observed for monomeric analogues.11 In addition, upon oxidation to 22+, this band disappears. NIR spectroelectrochemistry (NIR SEC) confirmed the reversibility of the oxidations (Figure S3, Supporting Information). This behavior is diagnostic of bands associated with intervalence charge transfer (IVCT) transitions, and we assign it as such. Below we discuss how the shape, intensity, and position of this IVCT transition confirm that 2•+ is more strongly coupled than 1•+. We first consider the shape of the IVCT band. Deconvolution of the IVCT band resulted in three Gaussian components (Figure 3). The low-energy transition with νmax at 6173 cm−1 is the most intense, and the two higher-energy Gaussian components decrease in intensity with increasing energy. It is not uncommon for IVCT bands to exhibit multiple components. In mixed-valence species in which the redox active unit is inorganic in nature, these features are commonly associated with multiple IVCT transitions (resulting from spin−orbit coupling) or with vibronic coupling.11 As 2 is an organic species, we rule out spin−orbit coupling, and instead we consider it most likely that these additional features are associated with a vibronic progression associated with the IVCT band.11,20 The spacing associated with these features observed for the IVCT band (1302 cm−1) is in close agreement with several of the modes found in the Raman spectra of 2•+ (Figure S7, Supporting Information). While assignment of these features to specific vibronic transitions is beyond the scope of this paper, we note that both the presence of strong vibronic features and asymmetry in the IVCT band are associated with the degree of strong electronic coupling exhibited by 2•+. In addition, the narrowness of the IVCT band is also an indicator of the degree of coupling, with the breadth decreasing with increasing coupling.12,21 The IVCT band of 2•+ is narrower (Δν = 1204 cm−1) than that of 1•+ (Δν

HAB = νmax /2 •+

(3) −1

•+

In 2 , νmax = 6173 cm , whereas, for 1 , νmax = 4660 cm−1. This difference corresponds to 757 cm−1 (0.17 V) stronger coupling in 2•+ than in 1•+ (Table 1), and our quantitative Table 1. Comparison of Electronic Coupling for 1•+, 2•+, and 3•+ ΔE1/2, V HAB, cm−1 HAB, eV Robin−Day classification

1•+a

2•+

3•+a

0.29 2330 0.29 III

0.35 3087 0.36 III

0.11 770b 0.095 II/III

a

From ref 11. bDetermined using the Hush formula for Class II systems.

analysis yields the result that 2•+ is 32% more strongly coupled than 1•+ despite a 53% (2.3 Å) increase in the distance between the redox units. As electronic coupling is expected to decrease with increasing distance between redox centers, there must be other factors besides distance affecting the coupling in these systems. We consider these factors next. 12697

dx.doi.org/10.1021/jp504887s | J. Phys. Chem. C 2014, 118, 12693−12699

The Journal of Physical Chemistry C

Article

changing from a biphenyl to a fluorene bridge).29 However, in many of these instances, the distance between the redox centers remains nearly constant, and thus does not exhibit the same behavior we observe for 2•+namely, an increase in coupling with increasing distance. However, at least one example (reported by Kochi and co-workers) does exist where electronic coupling was found to increase with increasing bridge length. Upon switching from a biphenylene to a stilbene bridge connecting two bis(alkoxy)benzene units, Kochi observed an increase of 188 cm−1 with a 2.5 Å increase in distance (this constitutes a general increase of 75 cm−1 Å−1).30 Kochi’s result is similar to ours; however, there are two points that deserve comment. First, as conceded by Kochi, their calculated coupling for the stilbene-bridged system may not be valid, as they employed the Hush equation. However, their system was assigned as near the Class II/III bordera region for which the validity of the Hush equation is suspect.30 Because both 1 and 2 are class III systems, our determination of the changes in HAB is clearer. Second, when increasing the distance between the directly connected units in 1 versus 2, we observed an increase in coupling of 757 cm−1 with an increase in distance of 2.3 Å (a general increase of 329 cm−1 Å−1). Thus, while both our results and Kochi’s indicate that introduction of ethylene between phenyl rings results in larger coupling, this effect is nearly 4.5 times stronger in our case. This result is, perhaps, expected as the stilbene-bridged system had larger distances between redox active units and retained phenyl− phenyl junctions. Both of these factors limit the ultimate extent of coupling that could be achieved by the increased planarity resulting from the addition of the ethylene moiety. In either case, both our results and those of Kochi support the hypothesis that the increased planarity introduced by the ethylene bridge leads to an increase in electronic coupling, despite an increase in distance between the redox centers.

The first factor that must be considered are any changes (between 1 and 2) in energy of the diabatic states of the molecules. Between these two molecules, both the bridge and the alkoxy units have changed. Both of these changes could affect the energy levels associated with the redox centers and the brdge. Because the degree of electronic coupling is inversely related to the energy gap between the redox units and the bridge, changes to these energetics would, in turn, affect the electronic coupling in the complexes. Given that we oxidize the molecules to generate the mixed valence state, it is the energy alignment of the HOMOs of the redox centers and the HOMO of the bridge that we must consider. The electrochemical data indicate that 2 is more easily oxidized (by 0.11 V) than either 1 or 3, indicating that the HOMO of the hexoxy-substituted donor unit in 2 is higher in energy than the methoxy-substituted donor unit in 1 or 3. This conclusion is arrived at by averaging the energies of the two redox potentials and taking this as an estimate of the diabatic energy (the energy in the absence of electrochemical splitting). This is an approach that has been used in the past for mixed valence systems.23 At the same time, the HOMO of ethene is lower than that of benzene, as measured by ionization potential,24 and so we also expect the HOMO of the ethylene bridge to lie lower in energy than the HOMO of a phenyl bridge. Thus, both of the structural differences in moving from 1 to 2 would be expected to increase the energy gap between the redox centers and the bridgea change that is also manifest in the blue shift of the UV−vis band associated with redox-site-to-bridge charge transfer (vide supra). As we just noted, the increase in this energy gap can only decrease the electronic coupling of 2•+ with respect to 1•+. We can therefore conclude that the energetic differences between 1 and 2 cannot account for the greater coupling observed for 2•+, with respect to 1•+. Instead of energetics, we suggest that the most important difference between 1 and 2 is the degree of steric hindrance. Specifically, the increase in distance results in a relief of steric hindrance in the molecule (originating from the spatial overlap of the phenyl rings), and this relief results in an increase in the planarity of 2•+ with respect to 1•+. Density functional theory (DFT) calculations support this conclusion by showing that the lowest-energy structures of 1 and 2 have a dihedral angle of 45° and 0° between the two aromatic rings, respectively (Figures S5 and S6, Supporting Information). This angle for 1 is in agreement with previously published crytallographic data that gave a dihedral engle of 69° (39° for the mixed-valence complex).11 These results also agree well with previously published computational studies.25,26 Here, we focus on the neutral compounds, as calculation is simpler than for the openshell mixed-valence state, and working with the neutral molecules provides the benefit of isolating the steric effects for the electronic coupling effects that are present in the mixedvalence state. The practical result of this analysis is that the rotational barriers to planarization, for 1 and 2, are calculated to be 8.8 and 0 kcal/mol, respectively. As the planar structure will give rise to the largest electronic coupling,27,28 there is a barrier to electronic coupling in 1•+ that is not present in 2•+. It is the absence of the additional barrier to electronic coupling that gives rise to the stronger coupling in 2•+, when compared to 1•+. There are, of course, examples in the literature where increasing the planarity of the bridge also increases the electronic coupling between redox centers (for example,

■

CONCLUSIONS

The results presented in this manuscript (considered with previous results)11 seem to demonstrate a general principle: electronic coupling between redox active units can be increased by inserting an ethylene bridge between two directly joined phenyl rings, and this increase in coupling is realized despite an increase in total distance between the redox sites. Thus, when weighing the need to maximize electronic coupling (which requires close proximity between redox sites), while minimizing steric hindrance (which requires distance between redox active units), there is a maximal compromise for which strong electronic coupling is possible, yet steric hindrance is minimal. In the case of bis(alkoxy)benzene redox active units (models for PPP and PPV) this maximal compromise appears to be realized by an ethylene bridge between the units, which provides 32% and 400% stronger coupling than for molecules with no bridge (1) or a phenyl bridge (3), respectively. This result quantitatively highlights the general utility of ethylene bridges for the design of future polymers, where conjugation length and packing are critical concerns. In addition, this result constitutes the first report of clearly defined ethylene-bridged mixed valency and highlights the advantages of ethylenes in mediating electron mobility within any molecule designed to carry electrons. 12698

dx.doi.org/10.1021/jp504887s | J. Phys. Chem. C 2014, 118, 12693−12699

The Journal of Physical Chemistry C

■

Article

Transitions in the Mixed-Valence Monocations of Bis(Triarylamines) Linked with Vinylene and Phenylene-Vinylene Bridges. J. Am. Chem. Soc. 2005, 127, 16900−16911. (13) Barlow, S.; Risko, C.; Coropceanu, V.; Tucker, N. M.; Jones, S. C.; Levi, Z.; Khrustalev, V. N.; Antipin, M. Y.; Kinnibrugh, T. L.; Timofeeva, T.; et al. A Mixed-Valence Bis(Diarylamino)Stilbene: Crystal Structure and Comparison of Electronic Coupling with Biphenyl and Tolane Analogues. Chem. Commun. 2005, 377, 764−766. (14) Jørgensen, M.; Krebs, F. C. Stepwise Unidirectional Synthesis of Oligo Phenylene Vinylenes with a Series of Monomers. Use in Plastic Solar Cells. J. Org. Chem. 2005, 70, 6004−6017. (15) Dennington, R.; Keith, T.; Millam, J. Gaussview, Version 5; Semichem Inc.: Shawnee Mission, KS, 2009. (16) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A., et al. Gaussian 09; Gaussian Inc.: Wallingford, CT, 2009. (17) Richardson, D.; Taube, H. Mixed-Valence Molecules: Electronic Delocalization and Stabilization. Coord. Chem. Rev. 1984, 60, 107−129. (18) Sutton, J. E.; Sutton, P. M.; Taube, H. Determination of the Comproportionation Constant for a Weakly Coupled Mixed-Valence System by Titration of the Intervalence Transfer Band: U-(4,4′Bipyridal)-Bis(Pentaamineruthenium)5+. Inorg. Chem. 1979, 18, 1017− 1021. (19) D’Alessandro, D. M.; Keene, F. R. A Cautionary Warning on the Use of Electrochemical Measurements to Calculate Comproportionation Constants for Mixed-Valence Compounds. Dalton Trans. 2004, 3950−3954. (20) Nelsen, S. F.; Tran, H. Q.; Nagy, M. A. Comparison of V Values for Some Nitrogen and Metal-Centered Pi-Bridged Mixed-Valence Compounds. J. Am. Chem. Soc. 1998, 120, 298−304. (21) Brunschwig, B. S.; Creutz, C.; Sutin, N. Optical Transitions of Symmetrical Mixed-Valence Systems in the Class II-III Transition Regime. Chem. Soc. Rev. 2002, 31, 168−184. (22) Heckmann, A.; Lambert, C. Neutral Organic Mixed-Valence Compounds: Synthesis and All-Optical Evaluation of ElectronTransfer Parameters. J. Am. Chem. Soc. 2007, 129, 5515−5527. (23) Londergan, C. H.; Salsman, J. C.; Lear, B. J.; Kubiak, C. P. Observation and Dynamics of ″Mixed-Valence Isomers″ and a Thermodynamic Estimate of Electronic Coupling Parameters. Chem. Phys. 2006, 324, 57−62. (24) Lias, S. G. Ionization Energy Evaluation. In NIST Chemistry Webbook, NIST Standard Reference Database Number 69; Linstrom, P. J., Mallard, W. G., Eds.; National Institute of Standards and Technology: Gaithersburg, MD, 2011. (25) Lhost, O.; Brédas, J. L. Theoretical Study of Torsional Potentials in Trans-Stilbene and Substituted Trans-Stilbenes: Modeling Torsions in Poly(Paraphenylene Vinylene) and Derivatives. J. Chem. Phys. 1992, 96, 5279−5288. (26) Pan, J.-F.; Chua, S.-J.; Huang, W. Conformational Analysis on Biphenyls with Theortical Calculations: Modeling Torsions in Poly(Para-Phenylene)s with Side Chains. Thin Solid Films 2000, 363, 1−5. (27) Bagrets, A.; Arnold, A.; Evers, F. Conduction Properties of Bipyridinium-Functionalized Molecular Wires. J. Am. Chem. Soc. 2008, 130, 9013−9018. (28) Schlüter, A.-D. Synthesis of Poly(Para-Phenylene)s. In Handbook of Conducting Polymers, 2nd ed.; Skotheim, T. A., Elsenbaumer, R. L., Reynolds, J. R., Eds.; Marcel Dekker: New York, 1998; pp 209− 224. (29) Heckmann, A.; Lambert, C. Organic Mixed-Valence Compounds: A Playground for Electrons and Holes. Angew. Chem., Int. Ed. Engl. 2012, 51, 326−392. (30) Rosokha, S. V.; Sun, D.-L.; Kochi, J. K. Conformation, Distance, and Connectivity Effects on Intramolecular Electron Transfer between Phenylene-Bridged Aromatic Redox Centers. J. Phys. Chem. A 2002, 106, 2283−2292.

ASSOCIATED CONTENT

S Supporting Information *

Details on the synthesis and characterization of 2; cyclic voltammogram of 2; UV−vis and NIR SEC spectra; IR and Raman spectra of 2 and 2•+; and results of DFT calculations. This material is available free of charge via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*Tel.: 814-867-4625. Fax: 814-865-3285. E-mail: bul14@psu. edu. Author Contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was supported in part through instrumentation funded by the National Science Foundation through grant OCI−0821527.

■

REFERENCES

(1) Allen, G. C.; Hush, N. S. Intervalence-Transfer Absorption I. Qualitative Evidence for Intervalence-Transfer Absorption in Inorganic Systems in Solution and in the Solid State. Prog. Inorg. Chem. 1967, 8, 357−390. (2) Winkler, J. R.; Gray, H. B. Electron Transfer in RutheniumModified Proteins. Chem. Rev. 1992, 92, 369−379. (3) Moratti, S. C. The Chemistry and Uses of Polyphenylenevinylenes. In Handbook of Conducting Polymers, 2 ed.; Skotheim, T. A., Elsenbaumer, R. L., Reynolds, J. R., Eds.; Marcel Dekker: New York, 1998; pp 343−362. (4) Jin, X.-H.; Sheberla, D.; Shimon, L. J.; Bendikov, M. Highly Coplanar Very Long Oligo(Alkylfuran)S: A Conjugated System with Specific Head-to-Head Defect. J. Am. Chem. Soc. 2014, 136, 2592− 2601. (5) Maior, R. M. S.; Hinkelmann, K.; Eckert, H.; Wudl, F. Synthesis and Characterization of Two Regiochemically Defined Poly(Alkylbithiophenes): A Comparative Study. Macromolecules 1990, 23, 1268−1279. (6) Gill, R. E.; Malliaras, G. G.; Wildeman, J.; Hadziiannou, G. Tuning of Photo- and Electroluminescence in Alkylated Polythiophenes with Well-Defined Regioregularity. Adv. Mater. 1994, 6, 132− 135. (7) Kanal, I. Y.; Owens, S. G.; Bechtel, J. S.; Hutchison, R. Efficient Computational Screening of Organic Polymer Photovoltaics. J. Phys. Chem. Lett. 2013, 4, 1613−1623. (8) O’Boyle, N. M.; Campbell, C. M.; Hutchison, G. R. Computational Design and Selection of Optimal Organic Photovoltaic Materials. J. Phys. Chem. C 2011, 115, 16200−16210. (9) Robin, A. M. B.; Day, P. Mixed Valence Chemistry - a Survey and Classification. Adv. Inorg. Chem. Radiochem. 1967, 10, 238−422. (10) Norris, B. N.; Zhang, S.; Campbell, C. M.; Auletta, J. T.; CalvoMarzal, P.; Hutchison, G. R.; Meyer, T. Y. Sequence Matters: Modulating Electronic and Optical Properties of Conjugated Oligomers Via Tailored Sequence. Macromolecules 2013, 46, 1384− 1392. (11) Lindeman, S. V.; Rosokha, S. V.; Sun, D.; Kochi, J. K. X-Ray Structure Analysis and the Intervalent Electron Transfer in Organic Mixed-Valence Crystals with Bridged Aromatic Cation Radicals. J. Am. Chem. Soc. 2002, 124, 843−855. (12) Barlow, S.; Risko, C.; Chung, S.-J.; Tucker, N. M.; Coropceanu, V.; Jones, S. C.; Levi, Z.; Brédas, J.-L.; Marder, S. R. Intervalence 12699

dx.doi.org/10.1021/jp504887s | J. Phys. Chem. C 2014, 118, 12693−12699